It is always possible to increase a rate
constant and accelerate a reaction by increasing the temperature.
For reactions with an activation energy of 12 to 13 kcal, at temperatures
around 2980K the rate constant doubles
with every 100 rise in temperature. (Can
you prove this?)
But as we have seen with NH3,
there can be difficulties: The reverse reaction may be accelerated
faster than the forward reaction, so that fewer products are obtained.
The products or reactants may be unstable at elevated temperatures,
or in special applications the surroundings may preclude the use
of higher temperatures.
For example, one cannot light a match
to burn glucose in the human body; this reaction must be carried
out at approximately 98.60F. It is for
such reactions that catalysis
In general, catalysts
lower the activation barrier for a reaction (Ea)
thereby making the rate
constants larger and the reactions faster. This is represented
schematically by the drawing opposite.
means finding an alternative pathway or mechanism for the reaction,
in which the intermediate states (activated
complexes) at all times are at a lower energy.
Both the forward and the reverse reactions
are speeded up by a catalyst, since lowering the forward Ea
necessitates lowering the reverse Ea
by the same amount. A catalyst has no effect on Keq
or on the ultimate equilibrium
conditions for a reaction; it only provides a way in which a spontaneous
but slow reaction can arrive at equilibrium faster.
If the reaction is not already thermodynamically
spontaneous, a cataIyst will be of no use. Thermodynamics does not
tell a chemist how to find a catalyst for a given reaction, but
it does tell him when it is, or is not, worth his time to look for